reactivity .doc

 American University of the Middle East REACTIVITY OF DIFFERENT METALS EXOTHERMIC ENDOTHERMIC Rxn Submitted to Dr . ABDULKADER BAROUDI Prepared by Fajer Alhaddad 13963 Laila Bekhtyari 18903 Ether Alzuwawi 17618 Amna Alramzi 17621 Taiba Alsaad 16757 Mona Alqhtani 17053 Ameena Al-Haid 15282 April 20 2014 table of contents abstract i table of contents ii 1 INTRODUCTION 1 1.1 reactivity 1 1.1.1 magnesium 1 1.1.2 hydrochloric 2 1.1.3 reaction HCL 2 1.1.4 reaction with H2so4 2 1.2 iron 3 1.2.1 structure of iron 3 1.2.2 reactivity of iron 4 1.2.3 iron with HCL 5 1.3 iron with h2so4 5 1.3.1 experimental 6 1.3.2 aim & procedure 6 1.3.resaults 6 1.3.4 refrance 7 REACTIVITY OF DIFFERENT METALS EXOTHERMIC & ENDOTHERMIC Rxn Fe & Mg Abstract Reactivity in chemistry refers to the chemical reactions of a single substance, the chemical reactions of two or more substances that interact with each other the systematic study of sets of reactions of these two kinds methodology that applies to the study of reactivity of chemicals of all kinds, experimental methods that are used to observe these processes theories to predict and to account for these processes. The chemical reactivity of a single substance (reactant) covers its behavior in which it: Decomposes Forms new substances by addition of atoms The reactivity of a metal is determined by how tightly bound the electrons of the element are. How easily and well a metal can replace other metals also determines the reactivity of the metal. The tighter the electrons are bound the less reactive metal. Introduction Reactivity of Metal Mg & Fe The activity series of metals is an empirical tool used to predict products in displacement reactions and reactivity of metals with water and acids in replacement reactions and ore extraction. It can be used to predict the products in similar reactions involving a different metal. ametal in the series, can displace any metal below it in the series, from the less reactive metal's oxide, chloride or sulphate or other compound. e.g. on heating the mixture of a metal and another metal oxide, such as magnesium powder and black copper(II) oxide, a very exothermic reaction occurs in a shower of sparks and white magnesium oxide is formed with brown bits of copper: Summary Theory From Study of some of the elements in the reactivity series and properties of each of magnesium and iron, magnesium turns out that more interactive and active than iron because it is preceded in reactivity series and activated quickly with air component of magnesium oxide also reacts with acids rapidly than iron . Magnesium was first extracted in 1808 by electrolysis. When magnesium oxide is very slightly soluble in water and forms magnesium hydroxide and the solution turns universal indicator solution If magnesium is heated in steam, the magnesium will burn with a bright white flame and the white powder magnesium oxide is formed and hydrogen gas. if the metal is at least as reactive as lead (see reactivity series list hydrochloric acid makes a metal chlorid salt ,and sulphuric acid makes a metal sulphate salt Mg(s) + 2HCl(aq) ==> MgCl2(aq) + H2(g) Mg (s) + H2SO4 (aq) ---> MgSO4 (aq) + H2 (g) In fact magnesium is so reactive, it will even burn in carbon dioxide, the products being white magnesium oxide powder and black specks of elemental carbon! The surface of iron goes dark grey-black when strongly heated in air/oxygen to form a tri-iron tetroxide. When steel wool is heated in a bunsen flame it burns with a shower of sparks - large surface area - increased rate of reaction - so even moderately reactive iron has its moments . Iron has no reaction with cold water to form hydrogen (rusting is a joint reaction with oxygen). When iron is heated in steam an iron oxide (unusual formula) and hydrogen are formed. This oxide is 'technically' diiron(III)iron(II) oxide but its sometimes called 'tri-iron tetroxide. Iron has a relative slow-moderate reaction with dilute hydrochloric acid forming the soluble pale green salt iron(II) chloride and hydrogen gas. Fe(s) + 2HCl(aq) ==> FeCl2(aq) + H2(g) It does not form iron(III) chloride, FeCl3, in this reaction, but it does form this other iron chloride compound when iron is heated in a stream of chlorine gas Iron has a slow reaction with dilute sulphuric acid forming the soluble pale green salt iron(II) sulphate and hydrogen gas. Fe(s) + H2SO4(aq) ==> FeSO4(aq) + H2(g) Experimental Reactivity series of metals|Reaction with hydrochloric acid and sulphuric acid Aim / Objective: To investigate the reaction of some metals: Mg,Fe, with hydrochloric acid and sulphuric acid. Apparatus/ Materials: metals of magnesium (Mg) turnings, zinc (Zn) granules, aluminum (Al) turnings, iron (Fe) filings, lead (Pb) foil, copper (Cu), 6 test-tubes, test-tube rack, dilute hydrochloric acid (HCl), dilute sulphuric acid (H2SO4), splints, Bunsen burner and spatula. Method / Procedure: Half fill a test- tube with dilute HCl. Add a spatula full of magnesium turnings to the acid, place cork stopper in the mouth of test-tube and observe for effervescence of gas. Record your observations. Heat gently under the Bunsen burner if no reaction is taking place or if it is too slow. Remove cork-stopper and place a lighted splint at the mouth of test-tube. Note the reaction. Repeat steps 1-4 with the remaining metals and dilute HCl. Then repeat the same procedure with dilute sulphric acid. Tabulate your results. Suggested Results:  METAL ACID SALT magnesium hydrochloric acid iron iron nitrate ethanoic acid sodium ethanoate calcium calcium sulphate copper nitric acid nitric acid iron nitrate sodium sodium chloride calcium ethanoic acid magnesium sulphuric acid  The surface of zinc goes white-yellow when strongly heated in air/oxygen to form zinc oxide (curiously ZnO is white when cold and yellow when hot due to an electron level effect). zinc + oxygen ==> zinc oxide 2Zn(s) + O2(g) ==> 2ZnO(s) Zinc oxide is insoluble with water.  Zinc has no reaction with cold water.  When the zinc is heated strongly in steam zinc oxide and hydrogen are formed. zinc + water ==> zinc oxide + hydrogen Zn(s) + H2O(g) ==> ZnO(s) + H2(g)  Zinc is quite reactive with dilute hydrochloric acid forming the colourless soluble salt zinc chloride and hydrogen gas. zinc + hydrochloric acid ==> zinc chloride + hydrogen Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g)  Zinc is quite reactive with dilute sulphuric acid forming the colourless soluble salt zinc sulphate and hydrogen gas. zinc + sulphuric acid ==> zinc sulphate + hydrogen Zn(s) + H2SO4(aq) ==> ZnSO4(aq) + H2(g) (this reaction is catalysed by adding a trace of copper sulphate solution which form a deposit on the zinc surface)  Zinc forms very little hydrogen with dilute nitric acid, though zinc nitrate is formed. This is because another reaction does occur in which the gas nitrogen monoxide (nitrogen(II) oxide, NO) is formed instead of hydrogen. The colourless nitrogen monoxide rapidly combines with oxygen in air to give the dangerous irritating brown gas nitrogen dioxide (nitrogen(IV) oxide, NO2). (i) zinc + nitric acid ==> zinc nitrate + hydrogen Zn(s) + 2HNO3(aq) ==> Zn(NO3)2(aq) + H2(g) which can occur in very dilute nitric acid  but has to compete with the reaction ... (ii) zinc + nitric acid ==> zinc nitrate + water + nitrogen(II) oxide [nitric oxide] 3Zn(s) + 8HNO3(aq) ==> 3Zn(NO3)2(aq) + 4H2O(l) + 2NO(g) and (ii) is rapidly followed rapidly by ... (iii)  nitrogen(II) oxide + oxygen ==> nitrogen(IV) oxide 2NO(g) + O2(g) ==> 2NO2(g) [nitric oxide ==> nitrogen dioxide] However with concentrated nitric acid, nitrogen dioxide is formed directly. (iv) zinc + nitric acid ==> zinc nitrate + water + nitrogen(IV) oxide Zn(s) + 4HNO3(aq) ==> Zn(NO3)2(aq) + 2H2O(l) + 2NO2(g) So, whatever concentration of nitric acid is used, you get a solution of zinc nitrate AND nasty brown fumes of nitrogen dioxide. Nitric acid is a strong oxidising agent and it is also NOT a reaction on which to base magnesium's position in the 'metal reactivity series' because of the complications.  Adding zinc granules to copper(II) sulphate solution, removes the blue colour of the copper(II) salt, leaving a colourless solution of zinc sulphate and a pinky-brown deposit of copper. zinc + copper sulphate ==> zinc sulphate + copper Zn(s) + CuSO4(aq) ==> ZnSO4(aq) + Cu(s) This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds.  Zinc can be extracted by reducing the hot metal oxide on heating with carbon  zinc oxide + carbon ==> zinc + carbon dioxide  2ZnO(s) + C(s) ==> 2Zn(s) + CO2(g)  A zinc coating (galvanising) is used to protect iron from rusting. The more reactive zinc oxidises 1st. Blocks of zinc attached to steel are used as 'sacrificial corrosion'.  Zinc was known and used in India and China before 1500 so it must have been extracted like copper or iron by carbon reduction of the oxide, sulphide or carbonate. 
Extraction of Zinc notes  The surface of aluminium goes white when strongly heated in air/oxygen to form white solid aluminium oxide. Theoretically its quite a reactive metal but an oxide layer is readily formed even at room temperature and this has quite an inhibiting effect on its reactivity.  Even when aluminium is scratched, the oxide layer rapidly reforms, which is why it appears to be less reactive than its position in the reactivity series of metals would predict but the oxide layer is so thin it is transparent,  so aluminium surfaces look metallic and not a white matt surface.  This property of aluminium makes it a useful metal for out-door purposes e.g. aluminium window frames, greenhouse frames. aluminium + oxygen ==> aluminium oxide 4Al(s) + 3O2(g) ==> 2Al2O3(s) Aluminium oxide is insoluble with water.  Under 'normal circumstances' in the school laboratory aluminium has virtually no reaction with water, not even when heated in steam due to a protective aluminium oxide layer of Al2O3. (see above) The metal chromium behaves chemically in the same way, forming a protective layer of chromium(III) oxide, Cr2O3, and hence its anti-corrosion properties when used in stainless steels and chromium plating. Although this again illustrates the 'under-reactivity' of aluminium, the Thermit Reaction shows its rightful place in the reactivity series of metals. The following is NOT needed for pre-university GCSE-AS-A2 etc. chemistry students as far as I'm aware, but maybe of interest to some students, because it illustrates what happens if you dig a little deeper into what appears to be a simple experimental situation! (1) If the surface of aluminium is treated with less reactive metal salt, it is still possible to get displacement reaction. Check this out by leaving a piece of aluminium foil in copper(II) sulphate solution and a patchy pink colour of copper metal slowly appears over many hours/days?. However, as a student teacher back in 1975, I did the experiment with a mercury salt (highly nerve toxic and now use banned in UK schools) and found all of the aluminium foil reacted when left in water overnight. The next morning, after the hydrogen had 'departed', there was nothing left but a soggy mass of hydrated aluminium hydroxide! The aluminium-mercury 'couple' enables the aluminium to displace the hydrogen from water even at room temperature. You get a similar 'speeding up' effect when copper(II) sulphate solution is added to a zinc-dilute sulphuric acid mixture. However, they are not as fast and exciting as the Thermit Reaction described below! which is legal for teachers to do with suitable health and safety precautions like using a transparent safety barrier and goggles and sending the class to the back of the room! (2) I am informed that water will react with molten aluminium because in the bulk of the liquid there is no oxygen. Thinking about, it does make sense if it is theoretically a reactive metal. Any traces of oxygen would be removed by the liquid aluminium forming Al2O3, leaving most of it un-oxidised. The reaction can then take place, and is very exothermically violent, forming the oxide/hydroxide and the flammable-explosive hydrogen gas. This is an important chemical health and safety issue encountered when dealing with metal extraction and foundry metal processes in industry well away from the relative 'small scale safety' of limited school industrial chemistry!  The Thermit reaction: However the true reactivity of aluminium can be spectacularly seen when its grey powder is mixed with brown iron(III) oxide powder. When the thermit mixture is ignited with a magnesium fuse (needed because of the very high activation energy!), it burns very exothermically in a shower of sparks to leave a red hot blob of molten=>solid iron and white aluminium oxide powder. Note the high temperature of the magnesium fuse flame is so high, the oxide layer (to the delight of all pupils) fails to inhibit the displacement reaction! yippee! (see above)  Equation and redox theory applied to the Thermite reaction aluminium + iron(III) oxide ==>  aluminium oxide + iron 2Al(s) + Fe2O3(s) ==> Al2O3(s) + 2Fe(s) The iron oxide is reduced to iron reduction is oxygen loss (Fe2O3 ==> Fe), aluminium is the reducing agent, gains the lost oxygen Fe3+ ions in Fe2O3 gain three electrons to form Fe atoms (electron gain is a more advanced definition of reduction), The aluminium is oxidised to aluminium oxide. oxidation is oxygen gain (Al ==> Al2O3), technically, iron oxide is the oxidising agent Al atoms lose three electrons to form Al3+ ions (in Al2O3) (electron loss is a more advanced definition of oxidation),  This is a typical displacement reaction by a more reactive metal displacing a less reactive metal from one of its compounds. In the blast furnace iron is displaced from iron oxide by using cheap carbon as the reducing agent. Aluminium is an expensive metal made by the costly process of electrolysis (Extraction of Aluminium), so the Thermit reaction would be a ridiculously expensive way of producing iron!  Slow reaction with dilute hydrochloric acid to form the colourless soluble salt aluminium chloride and hydrogen gas. (see above) aluminium + hydrochloric acid ==> aluminium chloride + hydrogen 2Al(s) + 6HCl(aq) ==> 2AlCl3(aq) + 3H2(g)  The reaction with dilute sulphuric acid is very slow to form colourless  aluminium sulphate and hydrogen. (see above) aluminium + sulphuric acid ==> aluminium sulphate + hydrogen 2Al(s) + 3H2SO4(aq) ==> Al2(SO4)3(aq) + 3H2(g)  If the surface of aluminium is treated with less reactive metal salt, it is still possible to get a displacement reaction. Check this out by leaving a piece of aluminium foil in copper(II) sulphate solution and a patchy pink colour of copper metal slowly appears over many hours/days? aluminium + copper(II) sulphate ==> aluminium sulphate + copper 2Al(s) + 3CuSO4(aq) ==> Al2(SO4)3(aq) + 3Cu(s)  Aluminium was first extracted in 1825 by electrolysis of its molten oxide Al2O3 (bauxite ore). The surface of iron goes dark grey-black when strongly heated in air/oxygen to form a tri-iron tetroxide. When steel wool is heated in a bunsen flame it burns with a shower of sparks - large surface area - increased rate of reaction - so even moderately reactive iron has its moments! iron + oxygen ==> iron oxide [iron tetroxide, diiron(III)iron(II) oxide] 3Fe(s) + 2O2(g) ==> Fe3O4(s) Iron oxide is insoluble with water. Iron has no reaction with cold water to form hydrogen (rusting is a joint reaction with oxygen). When iron is heated in steam an iron oxide (unusual formula) and hydrogen are formed. This oxide is 'technically' diiron(III)iron(II) oxide but its sometimes called 'tri-iron tetroxide'. iron + water (steam) ==> iron tetroxide + hydrogen 3Fe(s) + 4H2O(g) ==> Fe3O4(s) + 4H2(g) This is a reversible reaction - if you pass hydrogen over heated iron tetroxide it is reduced to iron and water is formed. iron tetroxide + hydrogen ==> iron + water (condenses) Fe3O4(s) + 4H2(g) ==> 3Fe(s) + 4H2O(g) Iron has a relative slow-moderate reaction with dilute hydrochloric acid forming the soluble pale green salt iron(II) chloride and hydrogen gas. iron + hydrochloric acid ==>  iron(II) chloride + hydrogen Fe(s) + 2HCl(aq) ==> FeCl2(aq) + H2(g) It does not form iron(III) chloride, FeCl3, in this reaction, but it does form this other iron chloride compound when iron is heated in a stream of chlorine gas (see salt preparation by direct synthesis note). Iron has a slow reaction with dilute sulphuric acid forming the soluble pale green salt iron(II) sulphate and hydrogen gas. iron + sulphuric acid ==> iron(II) sulphate + hydrogen Fe(s) + H2SO4(aq) ==> FeSO4(aq) + H2(g) Iron can be extracted by reducing the hot metal oxide on heating with carbon monoxide formed from carbon in the blast furnace e.g. iron(III) oxide + carbon monoxide ==> iron + carbon dioxide Fe2O3(s) + 3CO(g) ==> 2Fe(l-s) + 3CO2(g) iron tetroxide + carbon monoxide ==> iron + carbon dioxide Fe3O4(s) + 4CO(g) ==> 3Fe(l-s) + 4CO2(g) http://highschoolchemistryguide.com/reactivity-series-of-metalsreaction-with-hydrochloric-acid-and-sulphuric-acid-mg-zn-al-fe-pb-cu/ Magnesium burns vigorously with a bright white flame when http://www.slideshare.net/Arrehome/metals-reactivity-seriesh ttp://www.docbrown.info/page03/Reactivityb.htm http://www.docbrown.info/page03/Reactivity.htm http://www.slideshare.net/Arrehome/metals-reactivity-series http://www.physics-chemistry-class.com/reaction-iron-hydrocloric-acid.jpg